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Chapter 10 “Chemical Quantities”

Chapter 10 “Chemical Quantities”. Section 10.1 The Mole: A Measurement of Matter. OBJECTIVES: Describe methods of measuring the amount of something. Section 10.1 The Mole: A Measurement of Matter. OBJECTIVES: Define Avogadro’s number as it relates to a mole of a substance.

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Chapter 10 “Chemical Quantities”

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  1. Chapter 10“Chemical Quantities”

  2. Section 10.1The Mole: A Measurement of Matter • OBJECTIVES: • Describe methods of measuring the amount of something.

  3. Section 10.1The Mole: A Measurement of Matter • OBJECTIVES: • Define Avogadro’s number as it relates to a mole of a substance.

  4. Section 10.1The Mole: A Measurement of Matter • OBJECTIVES: • Distinguish between the atomic mass of an element and its molar mass.

  5. Section 10.1The Mole: A Measurement of Matter • OBJECTIVES: • Describe how the mass of a mole of a compound is calculated.

  6. What is a Mole? • You can measure mass, • or volume, • or you can count pieces. • We measure mass in grams. • We measure volume in liters. • We count pieces in MOLES.

  7. Moles (abbreviated: mol) • Defined as the number of carbon atoms in exactly 12 grams of carbon-12. • 1 mole is 6.02 x 1023 representative particles. • Treat it like a very large dozen • 6.02 x 1023 is called: Avogadro’s number.

  8. Representative particles • The smallest pieces of a substance: • For a molecular compound: it is the molecule. • For an ionic compound: it is the formula unit (made of ions). • For an element: it is the atom. • Remember the 7 diatomic elements (made of molecules)

  9. Types of questions • How many oxygen atoms in the following? • CaCO3 • Al2(SO4)3 • How many ions in the following? • CaCl2 • NaOH • Al2(SO4)3

  10. Some practice problems: • How many molecules of CO2 are there in 4.56 moles of CO2 ? • How many moles of water is 5.87 x 1022 molecules? • How many atoms of carbon are there in 1.23 moles of C6H12O6 ? • How many moles is 7.78 x 1024 formula units of MgCl2?

  11. Measuring Moles • Remember relative atomic mass? • The amu was one twelfth the mass of a carbon-12 atom. • Since the mole is the number of atoms in 12 grams of carbon-12, • the decimal number on the periodic table is also the mass of 1 mole of those atoms in grams.

  12. Gram Atomic Mass (gam) • Equals the mass of 1 mole of an element in grams • 12.01 grams of C has the same number of pieces as 1.008 grams of H and 55.85 grams of iron. • We can write this as: 12.01 g C = 1 mole C (this is also the molar mass) • We can count things by weighing them.

  13. Examples • How much would 2.34 moles of carbon weigh? • How many moles of magnesium is 24.31 g of Mg? • How many atoms of lithium is 1.00 g of Li? • How much would 3.45 x 1022 atoms of U weigh?

  14. What about compounds? • in 1 mole of H2O molecules there are two moles of H atoms and 1 mole of O atoms • To find the mass of one mole of a compound • determine the moles of the elements they have • Find out how much they would weigh (from the periodic table) • add them up

  15. Calculating Formula Mass Calculate the formula mass of magnesium carbonate, MgCO3. 84.32 g 24.31 g + 12.01 g + 3(16.00 g) =

  16. Section 10.2Mole-Mass and Mole-Volume Relationships • OBJECTIVES: • Describe how to convert the mass of a substance to the number of moles of a substance, and moles to mass.

  17. Section 10.2Mole-Mass and Mole-Volume Relationships • OBJECTIVES: • Identify the volume of a quantity of gas at STP.

  18. Molar Mass • Molar mass is the generic term for the mass of one mole of any substance (expressed in grams/mol) • The same as: 1) Gram Molecular Mass, 2) Gram Formula Mass, and 3) Gram Atomic Mass • just a much broader term.

  19. Examples • Calculate the molar mass of the following and tell what type it is: • Na2S • N2O4 • C • Ca(NO3)2 • C6H12O6 • (NH4)3PO4

  20. Molar Mass is… • The number of grams of 1 mole of atoms, ions, or molecules. • We can make conversion factors from these. • To change grams of a compound to moles of a compound.

  21. For example • How many moles is 5.69 g of NaOH?

  22. For example • How many moles is 5.69 g of NaOH?

  23. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles

  24. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH

  25. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g

  26. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g • 1 mole NaOH = 40.00 g

  27. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g • 1 mole NaOH = 40.00 g

  28. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g • 1 mole NaOH = 40.00 g

  29. The Mole-Volume Relationship • Many of the chemicals we deal with are gases. • They are difficult to weigh (or mass). • Need to know how many moles of gas we have. • Two things effect the volume of a gas • Temperature and pressure • We need to compare them at the same temperature and pressure.

  30. Standard Temperature and Pressure • 0ºC and 1 atm pressure • abbreviated “STP” • At STP 1 mole of any gas occupies 22.4 L • Called the molar volume • 1 mole = 22.4 L of any gas at STP

  31. Practice Examples • What is the volume of 4.59 mole of CO2 gas at STP? • How many moles is 5.67 L of O2at STP? • What is the volume of 8.8 g of CH4 gas at STP?

  32. Density of a gas • D = m / V • for a gas the units will be g / L • We can determine the density of any gas at STP if we know its formula. • To find the density we need: 1) mass and 2) volume. • If you assume you have 1 mole, then the mass is the molar mass • And, at STP the volume is 22.4 L.

  33. Practice Examples • Find the density of CO2 at STP. • Find the density of CH4 at STP.

  34. Another way: • Given the density, we can find the molar mass of the gas. • Again, pretend you have 1 mole at STP, so V = 22.4 L. • m = D x V • m is the mass of 1 mole, since you have 22.4 L of the stuff. • What is the molar mass of a gas with a density of 1.964 g/L? • How about a density of 2.86 g/L?

  35. Summary • These four items are all equal: a) 1 mole b) molar mass (in grams/mol) c) 6.02 x 1023 representative particles d) 22.4 L of gas at STP Thus, we can make conversion factors from them.

  36. Section 10.3Percent Composition and Chemical Formulas • OBJECTIVES: • Describe how to calculate the percent by mass of an element in a compound.

  37. Section 10.3Percent Composition and Chemical Formulas • OBJECTIVES: • Interpret an empirical formula.

  38. Section 10.3Percent Composition and Chemical Formulas • OBJECTIVES: • Distinguish between empirical and molecular formulas.

  39. Calculating Percent Composition of a Compound • Like all percent problems: Part whole • Find the mass of each of the components (the elements), • Next, divide by the total mass of the compound, then x 100 x 100 %

  40. Example • Calculate the percent composition of a compound that is 29.0 g of Ag with 4.30 g of S.

  41. Getting it from the formula • If we know the formula, assume you have 1 mole. • Then you know the mass of the pieces and the whole(these values come from the periodic table).

  42. Examples • Calculate the percent composition of C2H4? • How about Aluminum carbonate? • Sample Problem 10.10, p.307 • We can also use the percent as a conversion factor • Sample Problem page 308

  43. Formulas Empirical formula: the lowest whole number ratio of atoms in a compound. • molecular formula = (empirical formula)n [n = integer] • molecular formula = C6H6 = (CH)6 • empirical formula = CH Molecular formula: the true number of atoms of each element in the formula of a compound.

  44. Formulas(continued) Formulas for ionic compounds are ALWAYS empirical (lowest whole number ratio). Examples: NaCl MgCl2 Al2(SO4)3 K2CO3

  45. Formulas(continued) Formulas for molecular compoundsMIGHT be empirical (lowest whole number ratio). Molecular: H2O C6H12O6 C12H22O11 Empirical: H2O CH2O C12H22O11

  46. Calculating Empirical • Just find the lowest whole number ratio • C6H12O6 • CH4N • It is not just the ratio of atoms, it is also the ratio of moles of atoms. • In 1 mole of CO2 there is 1 mole of carbon and 2 moles of oxygen. • In one molecule of CO2 there is 1 atom of C and 2 atoms of O.

  47. Calculating Empirical • We can get a ratio from the percent composition. • Assume you have a 100 g. • The percentages become grams. • Convert grams to moles. • Find lowest whole number ratio by dividing by the smallest value.

  48. Example • Calculate the empirical formula of a compound composed of 38.67 % C, 16.22 % H, and 45.11 %N. • Assume 100 g so • 38.67 g C x 1mol C = 3.220 mole C 12.01 gC • 16.22 g H x 1mol H = 16.09 mole H 1.01 gH • 45.11 g N x 1mol N = 3.219 mole N 14.01 gN Now divide each value by the smallest value

  49. Example • The ratio is 3.220 mol C = 1 mol C 3.219 mol N 1 mol N • The ratio is 16.09 mol H = 5 mol H 3.219 mol N 1 mol N = C1H5N1 • A compound is 43.64 % P and 56.36 % O. What is the empirical formula? • Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula?

  50. Empirical to molecular • Since the empirical formula is the lowest ratio, the actual molecule would weigh more. • By a whole number multiple. • Divide the actual molar mass by the empirical formula mass – you get a whole number to increase each coefficient in the empirical formula • Caffeine has a molar mass of 194 g. what is its molecular formula?

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